Definition
• The study of the interchange of
chemical and electrical energy.
• Involve oxidation-reduction reactions
Half-Reaction Method
• Used to balance redox rxns in aqueous
solution
• One of the rxns is the oxidation rxn
• Other is the reduction rxn
• The half reactions are balanced
separately and then added
• Differs in acidic vs. basic solution
In Acidic Solution
1. Separate the oxid. and red. parts of the rxn.
2. For each:
a. Balance all elements but H and O
b. Balance O using H2O
c. Balance H using H+
d. Balance charge using electrons
3. If needed, mult. One or both half-rxns. by a
whole # to make electrons equal for both
4. Add half-rxns and cancel same on both sides
5. Make sure elements/charges are balanced
Acronym for Balancing in Acid
• Big White Hogs Eat Bagels and Cream Cheese
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B = balance all but H & O
W = add water for O’s
H = Add H
E = Add electrons
B = Make sure electrons match
CC = Combine and cancel
Example in Acid
• Balance the following equation in acid solution
using the half-reaction method.
Cu(s) + HNO3(aq) -> Cu2+(aq) + NO(g)
• Answer:
3Cu + 2HNO3 + 6H+ -> 3Cu2+ + 2NO + 4H2O
Cr2O72-(aq) + NO(g) -> Cr3+(aq) + NO3-(aq)
• Answer:
Cr2O72-(aq) + 2NO(g) + 6H+(aq) -> 2Cr3+(aq) + 2NO3(aq) + 3H2O(l)
In Basic Solution
• Follow the steps used in acidic solution
THEN…
• Add OH- to match H+ on same side. Add to
other side as well
• Form H2O by adding OH- and H+
• Cancel H2O’s on both sides
• Check for element/charge balance
Example in Base
• Balance the following equation in basic solution:
P4 -> PH3 + HPO32• Answer
P4 + 2H2O + 4OH- -> 2PH3 + 2HPO32-
Galvanic Cell Concept
• Separating the oxidation and reduction half-reactions
will make it so the energy in the electron transfer can
be harnessed.
• A salt bridge or porous disk is needed to connect the
half-reactions so ions can flow and electrons don’t
build up on one side of the reaction (one beaker)
• **VIDEO?
Galvanic Cell Definition
• Device which chemical energy
is changed to electrical energy.
• Oxidation occurs at the
ANODE
• Reduction occurs at the
CATHODE (cat gets fat =
cathode gains electrons a.k.a.
reduction)
– An ox and a red cat
(anode/oxidation,
reduction/cathode)
Electrodes
• If there is an element (not ion) in either
half-reaction, it is what that particular
electrode is made of ***comes up later!
• When all reactants/products are in
solution (aq) Pt or graphite can be used
Galvanic Cell Picture
(Parts)
Cell Potential/Electromotive Force
• (EMF) Represented by E°cell
• Unit = volt (V) = 1 joule/coulomb
• Measured with a voltmeter (not completely
accurate b/c of heat). A potentiometer is used
instead where the maximum cell potential can
be measured.
Standard Reduction Potentials
• If we can find the potential for each halfreaction (Table 18.1 pg. 829), we can
determine the cell potential (E°cell)
• Half-reaction manipulations (DO NOT
MANIPULATE VOLTAGE):
– One must be reversed (oxidation)…can reverse E
so you have -E = -voltage…
– Electrons lost must = electrons gained, so
multiplication of reaction may be needed (DO NOT
MULTIPLY VOLTAGE BY THIS NUMBER!)
• EQUATION (MEMORIZE)
E°cell = E°(cathode) - E°(anode)
Table 18.1
• Better oxidizing agents: easily reduced, LEFT
side rxn. = largest, most positive standard
reduction potential
• Better reducing agents: easily oxidized, RIGHT
side = most negative standard reduction
potential (aka most positive standard oxidizing
potential)
• Example: Which is best reducing agent Cu+, F-,
H-, H2O, I2, K (find in products)
• Answer: K (-2.92) > H- (-2.23) > Cu+ (0.16) > I2
(1.20) > H2O (1.23) > F- (2.92)
Galvanic Cell Example
• Calculate the emf values (E°cell) for the
following
Mg(s) + 2H+(aq) -> Mg2+(aq) +H2(g)
• Answer: E°cell = +2.37V
In Galvanic Cells…
• When E°cell is positive, the reaction will
run spontaneously. (last slide)
• If negative, it will run in the opposite
direction (will NOT run spontaneously
as written).
Cell Potential, Electrical Work,
and Free Energy
• Spontaneous IF:
– Positive cell potential
– Negative Gibbs Free Energy (we will learn
more about Gibbs Free Energy in a later
chapter)
Line Notation
• Not required for AP Exam
• A double vertical line separates the anode on
left and cathode on right
– Represents a salt bridge or porous disk
• A single vertical line separates different phases
• Ex:
anode Cd -> Cd2+ + 2ecathode Hg2+ + 2e- -> Hg
Line notation: Cd(s) I Cd2+(aq) II Hg2+(aq) I Hg
Galvanic Cell: Complete Description
In half-reaction descriptions…FOUR items needed:
1. The cell potential - positive when E°cell =
E°(cathode) - E°(anode) and the balanced cell
reaction
2. The direction of electron flow, obtained by inspecting
the half-reactions and using the direction that gives a
positive E°cell.
3. Designation of the anode and cathode.
4. The nature of each electrode and the ions present in
each compartment. A chemically inert conductor is
required if none of the substances participating in
the half-reaction is a conducting solid.
Example: Complete Description
•
Describe a galvanic cell based on the two halfreactions below.
Cu2+ + 2e- -> Cu E° = 0.34 V
Cr2O72- + 14H+ + 6e- -> 2Cr3+ + 7H2O E° = 1.33V
1. Balanced cell rxn: 3Cu(s) + Cr2O72-(aq) + 14H+(aq) ->
3Cu2+(aq) + 2Cr3+(aq) + 7H2O(l)
E°cell = 0.99V (needs to be positive)
2. 1.33 (cathode) -0.34 (anode) means Cu needs to be
reversed. Cu will be giving off e- which will travel
from Cu (anode) to cathode (platinum electrode).
Continued…
3. Anode (copper metal electrode), cathode
(platinum electrode)
4. The copper metal electrode (anode) will be
in the Cu/Cu2+ compartment while the
platinum electrode (cathode) will be in the
Cr2O72-/Cr3+ compartment
Changing Concentration
• In chemical reactions…
• Standard conditions = 1M for all
• When reactants are >1M, it will increase
product concentration and will increase the cell
potential
• When products are >1M, it will decrease the
product concentration (oppose forward
reaction), decreases cell potential
• Mg(s) + 2H+(aq) -> Mg2+(aq) +H2(g)
Concentration Cell
• A galvanic cell that has the same component on
each side but at different concentrations
• Causes a cell potential
• Voltages are typically small
• Ex: a cell has on its left side a 0.20 M Cu2+
solution and a 0.050 M Cu2+ solution on the
right side
Reaction Quotient
• Represented by Q
• Q = concentration of productsx
[reactants]y
• *Solids and liquids are not included…
• Multiply concentrations of products/reactants by
one another
• X and Y represent coefficients of each reactant
and product respectively
Reaction Quotient Example (13.5)
H2(g) + I2(g) <-> 2HI(g)
[H2]o = 0.81 M, [I2]o = 0.44 M, [HI]o = 0.58 M
Q = __[HI]2_ =
[H2]1[I2]1
(0.58)2
(0.81)(0.44)
= 0.94
Nernst Equation
• At 25°C…to find actual cell potential
• Ecell = E°cell - 0.0591 log(Q)
n
• n is number of moles of electrons transferred
• Q is the reaction quotient (see previous slide)
• E°cell is calculated using standard reduction
potentials (learned earlier this chapter)
Nernst Equation Example
• Calculate Ecell for a galvanic cell based on the
following half-reactions at 25C.
(Eq 1) FeO42- + 8H+ + 3e- -> Fe3+ + 4H2O
E° = +2.20 V
(Eq 2) O2 + 4H+ + 4e- -> 2H2O
E° = +1.23 V
[FeO42-] = 2.0 X 10-3 M, [Fe3+] = 1.0 X 10-3 M,
[O2] = 1.0 X 10-5 M, [H+] = 6.31 X 10-6 M
(see pg. 399 in study guide) E = 0.54 V
Ion-Selective Electrodes
• Cell potential is related to concentrations
• Electrodes can be used that are sensitive to
specific ions
• They measure concentrations of specific ions
which can have an effect on the cell potential
Batteries
• Groups of galvanic cells connected
which add together to give the total
battery potential
• Different types…
Lead Storage Battery
• Used in cars
• Lead = anode
• Lead coated with lead
dioxide = cathode
• Electrodes are dipped in
sulfuric acid
Dry Cell Batteries
• Calculators, electronic games, digital watches,
portable audio players, etc.
• Acid version, alkaline version, silver cell,
mercury cell, etc.
Fuel Cells
• Reactant constantly supplied to a
galvanic cell
• Reaction is:
2H2(g) + O2(g)  2H2O(l)
Corrosion
• Metals corrode because they oxidize easily
(spontaneously)
– The metals’ reduction potentials are less positive
than oxygen gas
– When half reactions are combined, there is a
positive E° value
– Speed of oxidation varies
– Some metals form a thin protective oxide coating,
preventing further erosion
Iron Corrosion
• Steel corrosion must be prevented b/c it makes
the backbone of most of our buildings, bridges,
and cars
• Chemical composition of steel is not uniform
– Places where iron is easily oxidized/reduced
• Ions and electrons migrate via moisture on
steel’s surface, causing rust
• Moisture acts as a salt bridge between the
anode/cathode regions
Corrosion Prevention
• A coating (paint/metal plating) is applied to
protect the metal from oxygen and moisture
– Chromium/tin used to plate steel with an oxide
coating
– Zinc used in galvanizing steel forms an oxidecarbonate coating. Oxygen reacts with the zinc
(sacrificial coating) instead of the metal underneath
• Alloying: covers metals with a thin layer of
stainless steel or other desirable alloy
More Corrosion Protection…
• Cathodic protection: for buried tanks/pipes
attaches an active metal (like Mg) loses
electrons, keeping iron from oxidizing
– Ships’ hulls work this way: bars of titanium
attached to the steel hull keep the hull from
being oxidized
Electrolysis
• Electrolytic cells use electrical energy to
produce a chemical change (requires energy)
– Electrolysis forces a current through a cell to
produce a chemical change for which the cell
potential is negative
– Causes a nonspontaneous rxn to occur
• Charges batteries, produces aluminum metal,
plates objects with chrome
• Switches the anode/cathode, flows in the
opposite direction as the galvanic cell
Stoichiometry in Electrochemistry
• Can measure how much chemical change
occurs with the flow of a given current for a
specific amount of time
• We add current and time to the chemical
equation and can use the mole to mole ratio to
do problems
• Important Conversion Factor for Ampere (amp):
1 A = coulomb/second
• 1 Faraday: 1 mole e- = 96,486 C
Stoichiometry Example
• How many grams of copper can be
reduced by applying a 3.00 A current for
16.2 min to a solution containing Cu2+
ions?
• Answer: 0.96 g Cu
Plating Order
• Plating means depositing neutral metal
on the electrode by reducing the metal
ions in solution
• More positive (higher) E° value will
plate out quickest
• Ex: pg. 851
Electrolysis Used Commercially…
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Production of Aluminum
Electrorefining of Metals
Metal Plating
Electrolysis of Sodium Chloride
Production of Aluminum
• Aluminum used to not be able to be produced
by humans (very expensive)
• Hall-Heroult process uses molten cryolite
(Na3AlF6) as a solvent for aluminum oxide
• Made aluminum easy (cheap) to produce
• Production of Al uses 5% of the total electricity
used in the U.S.
Electrorefining of Metals
• Purifies metals like Cu, Fe, Zn, Ag, Au, Pt
Metal Plating
• Metals that corrode easily can be plated
with a thin coating of a metal that is less
likely to corrode
• Ex: “tin” cans = steel cans with thin tin
coating, bumpers in cars are steel
coated in chrome
• Object needing coating serves as the
cathode in ions of the coating metal
Electrolysis of Sodium Chloride
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How sodium metal is primarily produced
NaCl mixed with CaCl2 to lower mp
Downs cell used for electrolysis
Pg. 856